Transcript Chapter 3: Elements, Compounds and the Periodic Table

Chapter 3:
Elements, Compounds,
and the Periodic Table
Chemistry: The Molecular Nature
of Matter, 6E
Jespersen/Brady/Hyslop
Discovery of Subatomic Particles
 Late 1800s & early 1900s
 Cathode ray tube experiments showed that
atoms are made up of subatomic particles
 Discovered negatively charged particles
moving from
 Cathode – negative electrode to
 Anode – positive electrode
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Discovery of Electron
JJ Thomson (1897)
 Modified cathode ray tube
 Made quantitative
measurements on
cathode rays
 Discovered negatively
charged particles
 Electrons (e)
 Determined charge to mass ratio (e/m) of these
particles
 e/m = 1.76 x 108 coulombs/gram
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Millikan Oil Drop Experiment
 Determining charge on Electron
 Calculated charge on electron
 e = 1.60 x 1019 C
 Combined with Thomson’s experiment to get
mass of electron
 m = 9.09 x 1028 g
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Discovery of Atomic Nucleus
Rutherford’s Alpha Scattering Experiment




Most alpha () rays passed right through gold
A few were deflected off at an angle
1 in 8000 bounced back towards alpha ray source
Gave us current model of nuclear atom
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Discovery Of Proton
 Discovered in 1918 in Ernest Rutherford’s lab
 Detected using Mass Spectrometer
 Hydrogen had mass 1800x mass of electron
 Masses of other gases whole number multiples of
mass of hydrogen
Proton
 Smallest
positively
charged particle
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Rutherford’s Nuclear Atom
 Demonstrated that nucleus:
 has almost all of mass in atom
 has all of positive charge
 is located in very small volume at center of atom
 Very tiny, extremely dense core of atom
 Where protons (p+) &
neutrons (1n) are
located
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Atomic Structure
 Electrons (e)
 Very low mass
 Occupy most of atom’s space
 Balance of attractive & repulsive forces controls
atom size
 Attraction between protons (p+) & electrons
(e) holds electrons around nucleus
 Repulsion between electrons helps them spread
out over volume of atom
 In neutral atom
 Number of es must equal number of p+s
 Diameter of atom ~10,000 × diameter of nucleus
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Discovery of Neutron (1n)
 First postulated by Rutherford & coworkers
 Estimated number of positive charges on nucleus
based on experimental data
 Nuclear mass based on this number of protons
always far short of actual mass
 About ½ actual mass
 Therefore, must be another type of particle
 Has mass about same as proton
 Electrically neutral
 Discovered in 1932 by Chadwick
 Caused free neutron to be created
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Properties of Subatomic Particles
Nucleus (protons
+ neutrons)
 3 Kinds of subatomic
particles of principal
interest to Chemists
Electrons
Particle
Mass (g)
Electrical
Charge
Electron
9.109391028
1
Proton
1.672641024
+1
0
1e
1
1
1 H, 1 p
Neutron
1.674951024
0
1
0n
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Symbol
10
Atomic Notation
Atomic number (Z)
 Number of protons that atom has in nucleus
 Unique to each type of element
 Element is substance whose atoms all contain
identical number of protons
 Z = # protons
Isotopes
 Atoms of same element with different masses
 Same number of protons (11 p )
 Different number of neutrons (10n )
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Atomic Notation
Isotope Mass number (A)
 A = (# protons) + (# neutrons)
 A=Z+N
 For charge neutrality, number of electrons &
protons must be equal
Atomic Symbols
 Summarize information about subatomic particles
 Every isotope defined by 2 numbers Z & A
A
 Symbolized by X
Z
Ex. What is the atomic symbol for helium?
He has 2
e –,
2n&2
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p+
Z = 2, A = 4
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2 He
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Isotopes
 Most elements are mixtures of 2 or more stable
isotopes
 Each isotope has slightly different mass
 Chemically, isotopes have virtually identical chemical
properties
 Relative proportions of different isotopes are
essentially constant
 Isotopes distinguished by mass number (A):
Ex.
 3 isotopes of hydrogen (H)
 4 isotopes of iron (Fe)
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Example:
What is the isotopic symbol for Uranium235?
 Number of protons (p+) = 92
= number of electrons in neutral atom
 Number of neutrons (1n) = 143
 Atomic number (Z) = 92
 Mass number (A) = 92 + 143 = 235
 Chemical symbol = U
 Summary for uranium-235: 235
92
U
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Learning Check:
 Fill in the blanks:
symbol
neutrons
60Co
33
81Br
46
65
29 Cu
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protons
27
electrons
27
35
35
29
29
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Your Turn!
206
An atom of 82 Pb has ___ protons, ___
neutrons, and ___ electrons.
A. 82, 206, 124
B. 124, 206, 124
C. 124, 124, 124
D. 82, 124, 82
E. 82, 124, 124
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Carbon-12 Atomic Mass Scale
 Need uniform mass scale for atoms
Atomic mass units (symbol u)
 Based on carbon:
 1 atom of carbon-12 = 12 u (exactly)
 1 u = 1/12 mass 1 atom of carbon-12 (exactly)
Why was 12C selected?
 Common
 Most abundant isotope of carbon
 All atomic masses of all other elements ~ whole
numbers
 Lightest element, H, has mass ~1 u
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Calculating Atomic Mass
 Generally, elements are mixtures of isotopes
Ex. Hydrogen
Isotope
1H
2H
Mass
1.007825 u
2.0140 u
%Abundance
99.985
0.015
How do we define Atomic Mass?
 Average of masses of all stable isotopes of given
element
How do we calculate Average Atomic Mass?
 Weighted average.
 Use Isotopic Abundances & isotopic masses
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Learning Check
Naturally occurring magnesium is a mixture of 3
isotopes; 78.99% of the atoms are 24Mg (atomic mass,
23.9850 u), 10.00% of 25Mg (atomic mass, 24.9858 u),
and 11.01% of 26Mg (atomic mass, 25.9826 u). From
these data calculate the average atomic mass of
magnesium.
0.7899 * 23.9850 u = 18.946 u
0.1000 * 24.9858 u = 2.4986 u
0.1101 * 25.9826 u = 2.8607 u
Total mass of average atom =
24.3053 u rounds up to 24.31 u
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24Mg
25Mg
26Mg
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Your Turn!
A naturally occurring element consists of two
isotopes. The data on the isotopes:
isotope #1
68.5257 u
60.226%
isotope #2
70.9429 u
39.774%
Calculate the average atomic mass of this element.
A. 70.943 u
0.60226 * 68.5257 u = 41.270 u
B. 69.487 u
0.39774 * 70.9429 u = 28.217 u
C. 69.526 u
69.487 u
D. 69.981 u
E. 69.734 u
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Periodic Table
 Summarizes periodic properties of elements
Early Versions of Periodic Tables
 Arranged by increasing atomic mass
 Mendeleev (Russian) & Meyer (German) in 1869
 Noted repeating (periodic) properties
Modern Periodic Table
 Arranged by increasing atomic number (Z):
 Rows called periods
 Columns called groups or families
 Identified by numbers
 1 – 18 standard international
 1A – 8A longer columns & 1B – 8B shorter columns
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Modern Periodic Table
with group labels and chemical families identified
Actinides
Note: Placement of elements 58 – 71 and 90 – 103 saves space
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Representative/Main Group Elements
A groups—Longer columns
 Alkali Metals
 1A = first group
 Very reactive
 All Metals except for H
 Tend to form +1 ions
 React with oxygen
 Form compounds that dissolve in water
 Yield strongly caustic or alkaline solution (M2O)
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Representative/Main Group Elements
A groups—Longer columns
 Alkaline Earth Metals
 2A = second group
 Reactive
 Tend to form +2 ions
 Oxygen compounds are strongly alkaline (MO)
 Many are not water soluble
 Accumulate in ground
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Representative/Main Group Elements
A groups—Longer columns
 Halogens
 7A = next to last group on right
 Reactive
 Form diatomic molecules in elemental state
 2 gases
 1 liquid
 2 solids
 Form –1 ions with alkali metals—salts
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Representative/Main Group Elements
A groups—Longer columns
 Noble Gases
 8A = last group on right
 Inert—very unreactive
 Only heavier elements of group react & then very
limited
 Don’t form charged ions
 Monatomic gases
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Transition Elements
B groups—shorter columns
 All are metals
 In center of table
 Begin in fourth row
 Tend to form ions with several different charges
Ex.
 Fe2+ and Fe3+
 Cu+ and Cu2+
 Mn2+, Mn3+, Mn4+, Mn5+, Mn6+, Mn7+
Note: Last 3 columns all have 8B designation
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Inner Transition Elements
Lanthanide elements
 Elements 58 – 71
Actinide elements
 Elements 90 – 103
 At bottom of periodic table
 Tend to form +2 and +3 ions.
 All Actinides are radioactive
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Metals, Nonmetals, or Metalloids
 Elements break down into 3 broad categories
 Organized by regions of periodic table
Metals
 Left-hand side
 Sodium, lead, iron, gold
Nonmetals
 Upper right hand corner
 Oxygen, nitrogen, chlorine
Metalloids
 Diagonal line between metals & nonmetals
 Boron to astatine
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Metals, Nonmetals, or Metalloids
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Metals
 Most elements in periodic table
Properties
 Metallic luster
 Shine or reflect light
 Malleable
 Can be hammered or
rolled into thin sheets
 Ductile
 Can be drawn into wire
 Hardness
 Some hard – iron & chromium
 Some soft – sodium, lead, copper
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Properties of Metals
 Conduct heat & electricity
 Solids at Room Temperature
 Melting points (mp) > 25 °C
 Hg only liquid metal (mp = –39 °C)
 Tungsten (W)
(mp = 3400 °C)
 Highest known for metal
 Chemical reactivity
 Varies greatly
 Au, Pt
very unreactive
 Na, K
very reactive
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Nonmetals
 17 elements
 Upper right hand corner of periodic table
 Exist mostly as compounds rather than as pure
elements
 Many are Gases
 Monatomic (Noble)
He, Ne, Ar, Kr, Xe, Rn
 Diatomic
H2, O2, N2, F2, Cl2
 Some are Solids: I2, Se8, S8, P4, C
 3 forms of Carbon (graphite, coal, diamond)
 One is liquid: Br2
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Properties of Nonmetals
 Brittle
 Pulverize when struck
 Insulators
 Non-conductors of
electricity and heat
 Chemical reactivity
 Some inert
 Noble gases
 Some reactive
 F2, O2, H2
 React with metals to form ionic compounds
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 8 Elements
Metalloids
 Located on diagonal line between metals &
nonmetals
 B, Si, Ge, As, Sb, Te, Po, At
Properties
 Between metals & nonmetals
 Metallic shine
 Brittle like nonmetal
 Semiconductors
 Conduct electricity
 But not as well as metals
 Silicon (Si) & germanium (Ge)
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Your Turn!
Which of the following statements is correct?
A. Cu is a representative transition element
B. Na is an alkaline earth metal
C. Al is a semimetal in group IIIA
D. F is a representative halogen
E. None of these are correct
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Your Turn!
All of the following are characteristics of metals
except:
A. Malleable
B. Ductile
C. Lustrous
D. Good conductors of heat
E. Tend to gain electrons in chemical reactions
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Ions
Ions & Ionic Compounds
 Transfer of 1 or more electrons from 1 atom to
another
 Form electrically charged particles
Ionic compound
 Compound composed of ions
 Formed from metal & nonmetal
 Infinite array of alternating Na+ & Cl ions
Formula unit
 Smallest neutral unit of ionic compound
 Smallest whole-number ratio of ions
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Formation of Ionic Compounds
Metal + Non-metal  ionic compound
2Na(s) + Cl2(g)
 2NaCl(s)
Na + Cl
+
Na

+ Cl
NaCl(s)
e
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Cations
Ionic Compounds
 Positively charged ions
 Formed from metals
 Atoms lose electrons
Ex. Na has 11 e– & 11 p+
Na+ has 10 e– & 11 p+
Anions
 Negatively charged ions
 Formed from non-metals
 Atoms gain electrons
Ex. Cl has 17 e– & 17 p+
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Cl– has 16 e– & 17 p+
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Experimental Evidence for Ions
Electrical conductivity requires charge movement
Ionic compounds:
 Do not conduct electricity in solid state
 Do conduct electricity in liquid & aqueous states
where ions are free to move
Molecular compounds:
 Do not conduct electricity in any state
 Molecules are comprised of uncharged particles
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Ions of Representative Elements
 Can use periodic table to predict ion charges
 When we use North American numbering of
groups: Cation positive charge = group #
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Ions of Representative Elements
 Noble gases are especially stable
Nonmetals
 Negative () charge on anion = # spaces you
have to move to right to get to noble gas
 Expected charge on O is
 Move 2 spaces to right
N
O
F
Ne
 O2–
 What is expected charge on N?
 Move 3 spaces to right
 N3 –
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Rules For Writing Ionic Formulas
1. Cation given first in formula
2. Subscripts in formula must produce
electrically neutral formula unit
3. Subscripts must be smallest whole numbers
possible

Divide by 2 if all subscripts are even

May have to repeat several times
4. Charges on ions not included in finished
formula unit of substance

If no subscript, then 1 implied
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Determining Ionic Formulas
Ex. Formula of ionic compound formed when
magnesium reacts with oxygen
 Mg is group 2A
 Forms +2 ion or Mg2+
 O is group 6A
 Forms –2 ion or O2–
 To get electrically neutral particle need
 1:1 ratio of Mg2+ & O2–
 Formula: MgO
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Determining Ionic Formulas
“Criss-cross” rule
 Make magnitude of charge on one ion into
subscript for other
 When doing this, make sure that subscripts are
reduced to lowest whole number.
Ex. What is the formula of ionic compound
formed between aluminum & oxygen ions?
Al3+ O2–
Al2O3
46
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Your Turn!
Which of the following is the correct formula for
the formula unit composed of potassium and
oxygen ions?
A. KO
B. KO2
C. K2O
D. P2O3
E. K2O2
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Your Turn!
Which of the following is the correct formula for
the formula unit composed of Fe3+ and sulfide
ions?
A. FeS
B. Fe3S2
C. FeS3
D. Fe2S3
E. Fe4S6
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Cations of Transition Metals
Transition metals
 Center (shorter) region of periodic table
 Much less reactive than group 1A & 2A
 Still transfer electrons to nonmetals to form ionic
compounds
 # of electrons transferred less clear
 Form more than 1 positive ion
 Can form more than 1
compound with same non-metal
Ex. Fe + Cl
FeCl2 & FeCl3
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Cations of Post-transition Metals
Post-transition metals
 9 metals Ga, In, Sn, Tl, Pb, Bi, Uut, Uuq, Uub
 After transition metals & before metalloids
 2 very important ones – tin (Sn) & lead (Pb)
 Both have 2 possible oxidation states
 Both form 2 compounds with same nonmetal
Ex. Ionic compounds of tin & oxygen are
 SnO & SnO2
 Bismuth
 Only has +3 charge
 Bi3+
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Ions of Some Transition Metals &
Post-transition Metals
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Compounds with Polyatomic Ions
Binary compounds
 Compounds formed from 2 different elements
Polyatomic ions
 Ions composed of 2 or more atoms linked by
molecular bonds
 If ions are negative, they have too many electrons
 If ions are positive, they have too few electrons
 Formulas for ionic compounds containing
polyatomic ions
 Follow same rules as ionic compounds
 Polyatomic ions are expressed in parentheses
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Table 3.4 Polyatomic Ions
(Alternate Name in
parentheses)
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Learning Check
Ex. What is the formula of the ionic compound
formed between ammonium and phosphate
ions?
 Ammonium = NH4+
 Phosphate = PO43–
(NH4)+ (PO4)3–
(NH4)3PO4
Ex. Between strontium ion and nitrate ion?
 Strontium = Sr2+
 Nitrate = NO32–
Sr2+ (NO3)–
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Sr(NO3)2
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Nomenclature (Naming)
 IUPAC system to standardize name of chemical
compounds
 One system so that anyone can reconstruct
formula from name
 We will look at naming Ionic Compounds of
 Representative metals
 Transition metals
 Monatomic ions
 Polyatomic ions
 Hydrates
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Naming Ionic Compounds
Cations:
 Metal that forms only 1 positive ion
 Cation name = English name for metal
 Na+
sodium
 Ca2+
calcium
 Metal that forms more than 1 positive ion
 Use Stock System
 Cation name = English name followed by numerical
value of charge written as Roman numeral in
parentheses (no spaces)
 Transition metal
 Cr2+
chromium(II)
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Cr3+
chromium(III)
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Naming Ionic Compounds
Anions:
 Monatomic anions named by adding
“–ide” suffix to stem name for element
 Polyatomic ions use names in Table 3.5
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Learning Check: Name The
Following
 K2O
potassium oxide
 NH4ClO3
ammonium chlorate
 Mg(C2H3O2)2
magnesium acetate
 Cr2O3
chromium(III) oxide
 ZnBr2
zinc bromide
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Learning Check: Determine The
Formula
 Calcium hydroxide
 Ca(OH)2
 Manganese(II) bromide
 MnBr2
 Ammonium phosphate
 (NH4)3PO4
 Mercury(I) nitride
 (Hg2)3N2
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Your Turn!
Which is the correct name for Cu2S?
A.
B.
C.
D.
E.
copper sulfide
copper(II) sulfide
copper(II) sulfate
copper(I) sulfide
copper(I) sulfite
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Your Turn!
Which is the correct formula for ammonium sulfite?
a) NH4SO3
b) (NH4)2SO3
c) (NH4)2SO4
d) NH4S
e) (NH4)2S
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Naming Hydrates
 Ionic compounds
 Crystals contain water molecules
 Fixed proportions relative to ionic substance
 Naming
 Name ionic compound
 Give number of water molecules in formula using
Greek prefixes
monoditritetrapenta-
=
=
=
=
=
1
2
3
4
5
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hexaheptaoctanonadeca-
=
=
=
=
=
6
7
8
9
10
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Learning Check: Naming Hydrates
 CaSO4 · 2H2O
 calcium sulfate dihydrate
 CoCl2 · 6H2O
 cobalt(II) chloride hexahydrate
 FeI3 · 3H2O
 iron(III) iodide trihydrate
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Your Turn!
What is the correct formula for copper(II) sulfate
pentahydrate?
A. CuSO4 · 6H2O
B. CuSO3 · 5H2O
C. CoSO4 · 4H2O
D. CoSO3 · 5H2O
E. CuSO4 · 5H2O
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Molecular Compounds
Molecules
 Electrically neutral particle
 Consists of two or more atoms
Chemical bonds
 Attractions that hold atoms together in molecules
 Arise from sharing electrons between 2 atoms
 Group of atoms that make up molecule behave as
single particle
Molecular formulas
 Describe composition of molecule
 Specify # of each type of atom present
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Molecules vs. Ionic Compounds
Molecules
 Discrete unit
 Water = 2 hydrogen atoms bonded to 1 oxygen atom
Ionic Compounds
 Ions packed as close as possible to each other
 Sodium chloride =
Each cation has 6
anions; each anion
has 6 cations
 No one ion “belongs”
to another
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Molecular Compounds
 Formed when nonmetals combine
 C + O2  CO2
2H2 + O2  2H2O
 Millions of compounds can form from a few nonmetals
 Organic chemistry & Biochemistry
 Deal with chemistry of carbon + H, N & O
 A few compounds have only 2 atoms
 Diatomics:
H2, O2, Cl2, HF, NO
 Most molecules are far more complex
 Sucrose (C12H22O11)
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urea (CON2H4)
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Hydrogen-containing Compounds
Nonmetal hydrides
 Molecule containing nonmetal + hydrogen
 Number of hydrogens that combine with nonmetal =
number of spaces from nonmetal to noble gas in
periodic table
N
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O
F Ne
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3-D Shapes of Molecules
 Space filling models
 Used to give shapes of simple nonmetal hydrides
 Blue = nitrogen
 Red = water
 Yellow = fluorine
 White = hydrogen
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Organic Compounds
 Carbon compounds
 Carbon + hydrogen, oxygen, & nitrogen
 Originally thought these compounds only came
from living organisms
 Now more general
Hydrocarbons
 Simplest organic compounds
 Contain only C & H
 Always have ratio of atoms CnH2n+2
 Named using prefix designating number of C atoms
 All have –ane suffix
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Table
3.8
Hydrocarbons Belonging to the
Alkane Series
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Alkanes
 Boiling point increases as number of carbon
atoms increases
 Space filling models of alkanes
 Black = carbon
 White = hydrogen
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Your Turn!
Which is the correct name for C4H10?
A. methane
B. ethane
C. propane
D. pentane
E. butane
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Other Hydrocarbons
Alkenes
 Hydrocarbons with two less H’s than alkanes
 CnH2n
 Name = number prefix + ene
Ex.
C2H4 = ethene (ethylene)
Alkynes
 Hydrocarbons with four fewer H’s than alkanes
 CnH2n – 2
 Name = number prefix + ene
Ex. C2H2 = ethyne (acetylene)
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Other Organic Compounds
 Hydrocarbons are basic building
blocks of organic chemistry
 Many other classes of
compounds derived from
them
Alcohols
 Replace H in alkane with -OH group
 Name = number prefix + anol
Ex. CH3OH = methanol (methyl alcohol)
C2H5OH = ethanol (ethyl alcohol)
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Your Turn!
What is the name of C4H9OH?
A. hexanol
B. propanol
C. pentanol
D. tetranol
E. butanol
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Writing Formulas for Organic Compounds
Molecular formula
 Indicates # of each type of atom in molecule
Ex. C2H6 for ethane or C3H8 for propane
 Order of atoms
 Carbon | Hydrogen | Other atoms alphabetically
Ex. sucrose is C12H22O11
Emphasize alcohol – write OH group last
 C2H5OH
Structural formula
 Indicate how carbon atoms are connected
 Ethane = CH3CH3
 Propane = CH3CH2CH3
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Your Turn!
Octane is a hydrocarbon with 8 C atoms that is
the major component of gasoline. What is the
correct molecular formula for octane?
A. C8H14
B. C8H16
C. C8H18
D. C8H17OH
E. C8H15OH
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Your Turn!
What is the correct structural formula for octane?
a) CH3CH2CH2CH2CH2CH2CH2CH3
b) CH3CH2CH2CH2CH2CH2CH3
c) C8H18
d) CH3CH2CH2CH2CH2CH2CH2CH2CH3
e) CH3CH2CH2CH2CH2CH2CH2CH2OH
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Nomenclature of Molecular Compounds
 Goal is a name that translates clearly into molecular
formula
Naming Binary Molecular Compounds
 Which 2 elements present?
 How many of each?
Format:
 First element in formula
 Use English name
 Second element
 Use stem & append suffix –ide
 Use Greek number prefixes to specify how many
atoms of each element
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Naming Binary Molecular Compounds
1. hydrogen chloride
1 H 1 Cl
2. phosphorous pentachloride
1 P 5Cl
3. triselenium dinitride
3 Se 2N
HCl
PCl5
Se3N2
 Mono always omitted on 1st element
 Often omitted on 2nd element unless more than one
combination of same 2 elements
Ex. Carbon monoxide
Carbon dioxide
CO
CO2
 When prefix ends in vowel similar to start of
element name, drop prefix vowel
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Learning Check: Name Each
Format:
 Number prefix + 1st element name
 Number prefix + stem + –ide for 2nd element
 AsF3
 HBr
=
=
arsenic trifluoride
 N2O4
 N2O5
 CO
=
=
=
dinitrogen tetroxide
 CO2
=
carbon dioxide
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hydrogen bromide
dinitrogen pentoxide
carbon monoxide
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Your Turn!
Which is the correct formula for nitrogen
triiodide?
A. N3I
B. NI3
C. NIO3
D. N(IO3)3
E. none of the above
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Your Turn!
Which is the correct name for P4O10?
A. phosphorus oxide
B. phosphorous decoxide
C. tetraphosphorus decoxide
D. tetraphosphorus oxide
E. decoxygen tetraphosphide
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Exceptions to Naming Binary Molecules
Binary compounds of nonmetals + hydrogen
 No prefixes to be used
 Get number of hydrogens for each nonmetal from
periodic table
 Hydrogen sulfide = H2S
 Hydrogen telluride = H2Te
Molecules with Common Names
 Some molecules have names that predate IUPAC
systematic names
 Water
H 2O
 Ammonia
NH3
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▪ Sucrose
C12H22O11
▪ Phosphine PH3
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Summary of Naming
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